AM Syllabus (2015): Chemistry
AM SYLLABUS (2015)
CHEMISTRY
AM 06
SYLLABUS
1
AM Syllabus (2015): Chemistry
Chemistry AM06
Syllabus
Paper 1 (3 hours) + Paper II (3 hours) + Paper III (3 hours)
Chemistry is not only a subject of academic study which has engaged scholars for centuries worldwide but its products
and processes have indeed enabled the very existence of societies and allowed their development. It is a central science
which informs several other disciplines including biology, medicine, materials science, agricultural science and others.
Chemical warfare and pollution of the environment as well as drug abuse, adulteration of food and similar examples of
misuse or improper use or disposal of chemicals often cast shadows on this scientific enterprise and generate mistrust
of chemicals by the general public.
The main purpose of this syllabus is to impart a proper understanding of the core principles and unifying ideas of
chemical science, supported and illustrated by a significant body of essential factual knowledge that should help
students acquire a meaningful experience of chemistry to serve them adequately in their future studies. The syllabus
also provides an opportunity for learners to become properly acquainted with this important science in order for them to
be able to better appraise both the opportunities and the challenges which chemistry and its products present to society
at large.
The syllabus for Chemistry at this level assumes and builds on knowledge of the subject at a level that is equivalent to
that covered by the Secondary Education Certificate syllabus.
The syllabus is presented in two columns, describing Subject Content and Learning Outcomes. This design is intended
to assist teachers and students determine the depth of treatment for each subject matter mentioned.
Teachers may of course go beyond the learning outcomes in class but the examination will test the outcomes as
described in this syllabus.
The course is intended to build both a theoretical knowledge base in chemistry that prepares students for further higher
studies in the subject as well as to develop basic practical skills. With respect to practical chemistry, students should
learn how to perform laboratory work in a manner which respects both their health and the environment; teachers
should evaluate the risks associated with practical work with a view to safeguard student safety and environmental
integrity.
The Examination
The examination consists of two written papers and a practical paper. Each of the papers is of three hours duration.
Paper I consists of six to ten compulsory structured questions and carries 40% of the total score. Students will write
their answers on the examination paper in the spaces provided.
Paper II consists of two sections each containing four extended response questions. Five questions are to be answered
from this paper, two questions from each section and any one other. Each question carries equal marks. This paper
carries 40% of the total score.
Paper III is a practical examination. This is an open book examination and candidates may use any printed material
which assists them in their work. This paper carries 20% of the total score.
Candidates may use an electronic calculator in all parts of the examination.
Familiarity with Periodic Table
Candidates are expected to be familiar with the structure of the Periodic Table and the group affiliation of each element
with atomic number from 1 to 30. However, a copy of the Periodic Table will be provided for Paper II of the
examination although not for Paper I. Relative atomic masses will always be given where necessary.
Mathematical skills
In order to understand certain concepts required by the syllabus, candidates will need to be able to:
(i)
Recognise and use expressions in decimal and standard form; use ratios, fractions and percentages and find
arithmetic means; make estimates of results; use an appropriate number of significant figures; use calculators
to find and use xy, x, 1/x, log10 x.
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AM Syllabus (2015): Chemistry
(ii)
Change the subject of an equation; substitute numerical values into algebraic equations using units for physical
quantities; use logarithms in relation to quantities which range over several orders of magnitude.
(iii)
Appreciate angles and shapes in regular two and three dimensional structures and to represent three
dimensional forms in two dimensions; understand the symmetry of two dimensional and three dimensional
shapes.
(iv)
Plot two variables from given data; understand that y = mx + c represents a linear relationship and be able to
determine the slope and intercept of a line; draw and use the slope of a tangent to a curve as a measure of rate
of change.
Language skills
Candidates are reminded of the importance of the use of good English in answering questions in the examination. In
essay-type answers and also where calculations are involved, orderly and detailed presentation will be rewarded.
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AM Syllabus (2015): Chemistry
Syllabus
Content
1
Atomic structure
Learning outcomes and comments
1.1
Students should be able to:
The fundamental particles: protons, electrons and
neutrons, their charges and relative masses.
describe the structure of an atom in terms of fundamental
particles using the Bohr model;
1.2
The nuclear structure of the atom. Proton (or
atomic) number and nucleon (or mass) number.
Isotopes and relative atomic masses. The carbon12 scale.
Students should be able to:
deduce the relative atomic mass of an element from the
isotopic masses and relative abundances;
describe the use of isotopes as tracers in mechanistic
studies using the example of of 18O labelling in
esterification;
1.3
Radioactivity: alpha and beta particles and
gamma rays. Half-life. Nuclear equations.
Students should be able to:
apply the concept of half life in determining the age of a
carbon-containing artefact;
determine the half-life of a radioactive isotope using an
exponential decay curve;
describe C-14 dating.
Positron emission will not be tested.
Calculations requiring use of the equation for radioactive
decay will not be set.
1.4
Introductory treatment of quantised energy levels
in atoms: evidence from atomic spectra; s, p and
d orbitals.
Students should be able to:
draw the electronic configuration of isolated atoms of
elements H to Kr using 1s, 2s, 2p etc notation and
electrons- in-boxes notation and applying the ‘building–
up’ (aufbau) principle.
Questions on interpretation of atomic spectra will not be
set although students are expected to explain how
transitions between energy levels can give rise to flame
colours.
1.5
Classification of elements into periods, groups
and ‘blocks’, including the first row d-block
elements.
1.6
Shapes of s and p orbitals. Hybridised orbitals.
Students should know the shape of hybrid orbitals sp, sp2
and sp3 and relate it to molecular geometry.
1.7
Ionisation energies and electron affinity and
relation to electronic configuration.
Students should be able to:
explain how ionisation energies vary across the second
and third period, both the general trends and irregularities
and down a group in the Periodic Table;
be able to explain why second and subsequent electron
affinities are endothermic.
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AM Syllabus (2015): Chemistry
1.8
The mass spectrometer. An elementary treatment
of mass spectrometry including the use of the
mass spectrometer to determine the relative
atomic mass and as a tool for determination of
molecular structure (see also section 2.10).
Students should be able to:
draw and label a diagram showing the principle of
operation of a magnetic sector mass spectrometer;
recognize and interpret a simple mass spectrum of an
atomic and a molecular sample, explaining the peak height
ratios for the mass spectra of Cl2 and Br2, and restricting
discussion to singly-charged peaks;
use the mass spectrum of an element to obtain the relative
atomic mass.
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AM Syllabus (2015): Chemistry
2
Electronic theory and chemical bonding
2.1
The ionic (electrovalent), covalent and coordinate (or dative) bond. Electronegativity.
Intermediate bonding: ion polarisation and bond
polarisation.
The concept of electron sharing and electron transfer
between atoms as complete or partial processes depending
on the electronegativity of the atomic species concerned
with reference to Pauling’s electronegativity scale. Dative
bonding exemplified by molecular compounds such as
CO, HNO3, Al2Cl6, NH3BF3 and ions H3O+ and NH4+.
2.2
Polar covalent bonds and electrical dipoles in
molecules (qualitative treatment only). Polar
covalent bonds which may, or may not, give rise
to molecules with a permanent dipole.
Vector representation of dipole moments to explain why
polar covalent bonds in molecules can lead to molecules
having a permanent dipole or a zero dipole: qualitative
treatment only. Use of symbols δ+ and δ- to represent
partial charges in bonds.
2.3
Nature of forces in bonding.
Comparison
between ionic and covalent bonding. Multiple
bonding.
Students should be able to describe:
Lewis structures drawn showing electrons as dots and
crosses;
covalent bonding in terms of overlap of orbitals, including
hybridised orbitals to produce σ and  bonds.
2.4
Delocalisation of electrons.
Evidence for delocalisation from bond lengths and
thermochemical data. Delocalised  bonds in, for
example, benzene, buta-1,3-diene, ozone and carbonate,
nitrate and alkanoate ions. Students should be familiar
with resonance structures (canonical forms) and resonance
hybrids (delocalised structures).
2.5
Metallic bonding
Electron sea model; use of model to account for
malleability, ductility and thermal and electrical
conductivity of metals. Students should be able to relate
the strength of metallic bonding to attraction between
positive ions and delocalised valence electrons and to
explain the variation of metallic bonding down a group.
2.6
Lattice structures of NaCl and CsCl as typical
ionic solids. Coordination numbers.
Students should be able to relate the ratio of cationic
radius to anionic radius to type of crystal lattice and to
understand the concept of the unit cell. Students should be
able to state the approximate (not exact) bond angles in
molecules and ions having both lone and bonding pairs
surrounding the central atom with up to six electron pairs
2.7
Shapes of molecules and ions and the valence
shell electron pair repulsion (VSEPR) theory.
2.8
Molecular
structures.
crystals
and
macromolecular
Students should be able to :
account for physical properties (melting point, boiling
point, electrical conductivity) in terms of the structure for
such substances as iodine, carbon dioxide, carbon
(diamond), carbon (graphite) and silicon(IV) oxide.
2.9
Intermolecular forces: hydrogen bonding, and
van der Waals forces (permanent dipolepermanent dipole and induced dipole-induced
dipole forces).
Students should be able to:
explain the effects of these forces on specific properties of
molecular compounds: e.g. boiling points of simple
hydrides in Groups 4, 5, 6 and 7; variation of the boiling
points along a homologous series of organic compounds;
formation of carboxylic acid dimers; the open structure of
ice leading to its density being lower than that of liquid
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AM Syllabus (2015): Chemistry
water;
distinguish between intermolecular and intramolecular
hydrogen bonding as in 4-nitrophenol and 2-nitrophenol
isomers.
2.10
Mass spectrometer as a method of determining
molecular structure: interpretation of simple mass
spectra involving only singly charged ions; the
molecular (parent) ion and fragment ions.
Use of the mass spectrum of a simple molecular substance
to obtain information about the molecular structure of the
substance. Knowledge of simple fragmentation only (no
knowledge of peak formation via rearrangements is
expected).
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AM Syllabus (2015): Chemistry
3
States of Matter
3.1
The ideal gas law.
Use of ideal gas law in the determination of relative
molecular mass for gases and volatile liquids: use of gas
syringe.
3.2
Dalton’s Law of Partial Pressures; diffusion and
effusion.
Graham’s Law of diffusion will not be examined.
3.3
Kinetic theory of gases; distribution of molecular
speeds.
Qualitative treatment only.
Students should be able to draw distribution diagrams
showing speed of gas molecules for different temperature
and to relate this to the concept of activation energy of
chemical reactions. See Section 10. The Zartmann
experiment will not be tested.
3.4
Real gases and deviation from ideal behaviour.
Van der Waals equation.
Basic assumptions of the ideal gas model. No numerical
questions will be set to test knowledge of van der Waals
equation.
3.5
Vapour pressure and saturation vapour pressure;
vaporisation and fusion in terms of the kinetic
molecular model.
Students should know about the relationship between
vapour pressure and temperature.
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AM Syllabus (2015): Chemistry
4
Quantity of matter
4.1
Moles of substance. Molar concentration in terms
of moles dm-3 or mol L-1; mass concentration in
terms of g dm-3 or g L-1. Volume measured in
dm3 or L and cm3 or mL.
4.2
Avogadro constant. Empirical and molecular
formulae.
4.3
Chemical equations, both full and ionic and the
use of these equations in calculations on reacting
substances in terms of amounts and
concentrations measured in moles and mol dm-3
respectively.
Students should be able to define the mole and to perform
chemical calculations in terms of molar quantities. Matter
described in terms of atoms, molecules, ions.
Students should be able to:
use in calculations the concept of limiting reagent;
calculate percentage yield and state why reactions do not
always produce 100% yield;
calculate the percentage purity of a substance from given
data.
4.4
Titrimetric analysis: acid-base, precipitation
(silver nitrate – chloride) redox and
complexometric (EDTA) reactions.
Students should be able to explain the experimental details
of titrimetry and be able to carry out a titrimetric analysis
(for more details see Section 14).
4.5
Back titrations
Students should be able to explain how to carry out back
titration involving acid/base and redox systems.
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AM Syllabus (2015): Chemistry
5
5.1
Energetics
Changes of energy accompanying phase changes
and chemical changes. Energy level (enthalpy)
diagrams. Joule (J) as the unit of energy.
5.2
Standard enthalpy change of reaction, formation,
atomisation, ionization, combustion, solution (or
dissolution),
neutralisation
and solvation
including hydration. Electron affinity. Bond
enthalpy terms and bond dissociation enthalpies;
lattice enthalpies; enthalpies of solution and
relation to lattice enthalpy and enthalpies of
hydration.
Students should be able to:
define the enthalpy changes listed in this section;
distinguish between first and second electron affinity for
dinegative ions;
explain the difference between bond enthalpy terms and
bond dissociation enthalpies;
use bond enthalpies to estimate the enthalpy change of a
reaction and explain why values obtained in this manner
are approximate.
5.3
Hess’s law and its use in simple calculations.
Born-Haber cycle.
Students should be able to:
state Hess’s Law and use it to calculate enthalpy changes
indirectly by construction of simple enthalpy cycles or
energy level diagrams from given thermochemical data;
construct Born-Haber cycles and carry out associated
calculations.
5.4
Comparison
between
theoretical
and
experimental values of enthalpy change and
interpretation in terms of structure and bonding.
Silver halides provide good examples of deviation
between theoretical and experimental lattice enthalpies
and interpretation of difference in terms of a degree of
covalency in this formally ionic structure.
The difference between expected and experimental values
of enthalpies of formation or combustion of benzene and
similar simple hydrocarbons provide evidence of
delocalised structures.
5.5
Calorimetry: determination of enthalpy change
of neutralization, solution and reaction. Calorific
value of fuels and food. Simple thermometric
acid-base titrations.
Students should be able to:
interpret temperature-time
curves to determine
temperature change and to use such data in the
relationship H = - mcT;
interconvert Joule and calorie as units of enthalpy change.
5.6
Concept of system and its surroundings. Entropy,
Gibbs free energy and spontaneity of chemical
change. Familiarity and use of the relationship
between free energy, enthalpy and entropy,
namely, G = H - TS. Kinetic versus
thermodynamic stability.
Students should be able to:
state that H alone is not able to fully predict spontaneous
change;
consider entropy as a measure of disorder of a system and
predict whether S is positive, negative or nearly zero in
simple changes such as state change, combustion,
dissolution, dimerization;
calculate S from absolute entropy values;
explain that processes are spontaneous when the free
energy change is negative and to calculate the temperature
at which processes start or cease to be spontaneous by use
of the equation G = H - TS.
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AM Syllabus (2015): Chemistry
6
Phase Equilibria
6.1
Phase changes for one-component systems: H2O,
CO2. Pressure-temperature phase diagrams and
boiling, melting and triple points.
Critical
temperature and supercritical fluids.
Students should be able to use phase diagrams to explain
sublimation and know about the use of supercritical fluids
as ‘green’ solvents which can replace toxic organic
solvents, e.g. in the food industry.
6.2
Two-component systems:
mixtures of two
miscible liquids and Raoult’s Law. Pressurecomposition
and
temperature-composition
diagrams.
Deviations from Raoult’s Law.
Azeotropic mixtures. Fractional distillation of
ideal and non-ideal mixtures.
Students should be able to:
give examples of binary mixtures which are
approximately ideal and which show both positive and
negative deviations from Raoult’s Law;
use Raoult’s law to calculate the vapour pressure of a
component above an ideal binary mixture;
employ Dalton’s law to determine the composition of the
vapour above an ideal binary mixture.
6.3
Immiscible liquids and steam distillation.
explain why a mixture of two immiscible liquids boil at a
temperature below the boiling point of both liquids;
describe the advantage of steam distillation over simple
distillation.
deduce an equation which relates the mass ratio of the
substances in the vapour phase to the vapour pressures of
the liquids and their relative molecular masses.
6.4
Osmosis and osmotic pressure: use of osmotic
pressure, , for the determination of relative
molecular mass from relation V = nRT where n
is the amount, in moles, of particles of solute.
Reverse osmosis.
Osmosis as an example of a colligative property of
solutions. Questions on the determination of partial degree
of dissociation/association will not be set.
Students should be able to relate reverse osmosis to local
desalination of sea water and compare it to other methods
of purification of water in terms of energy considerations.
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AM Syllabus (2015): Chemistry
7
7.1
Chemical Equilibrium
Concept of dynamic equilibrium. Characteristics
of an equilibrium mixture. Equilibrium constant
in terms of concentrations and of pressures: KC
and KP including units as appropriate.
Homogeneous and heterogeneous equilibria and
their associated equilibrium constants.
Students should be able to:
identify the main characteristic of the equilibrium state as
constancy of concentrations/pressures resulting from the
equality of the rates of the forward and reverse reactions;
deduce the extent of reaction from the magnitude of the
equilibrium constant;
write equilibrium constant expressions involving
condensed (liquid or solid) phases in heterogeneous
equilibria;
7.2
Degree of dissociation
write an expression for the degree of dissociation of
molecular covalent substances converting into smaller
molecules;
relate the degree of dissociation to KC and KP. Problems
requiring solving quadratic equations using the formula
will not be set.
7.3
Experimental methods of investigating chemical
equilibrium.
Students should be able to describe how the following
equilibrium reactions can be investigated practically: (a)
esterification reaction of a carboxylic acid and an alcohol;
(b) dissociation of hydrogen iodide into the elements
7.4
Le Chatelier’s Principle.
The effect of
concentration, temperature, and (where relevant)
pressure on: (a) wholly gaseous equilibria, (b)
solid-gas equilibria, (c) equilibria in the liquid
phase (e.g. esterification) or in solution. Effects
of presence of catalysts on equilibrium.
Students should be able to:
use the Haber Process for the manufacture of ammonia to
illustrate this principle and to explain how and why
compromise conditions are required. (Details of
manufacturing plant will not be tested);
describe and predict the qualitative effect of changes of
temperature or concentration or pressure on the
equilibrium position and on the value of the equilibrium
constant;
describe the effect of the introduction of an inert gas on
an equilibrium mixture at constant volume.
7.5
Distribution of a non-volatile solute between two
immiscible solvents: partition constant; solvent
extraction.
Students should be able to solve numerical questions on
partition constant Kd related to solvent extraction.
Discussion of association and dissociation effects on
partition equilibrium are not required.
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AM Syllabus (2015): Chemistry
8
8.1
Ionic Equilibria
Acids and bases: Arrhenius, Bronsted-Lowry
and Lewis definitions; conjugate pairs; proticity
or ‘basicity’ of acids and ‘acidity’ of bases (e.g.
sulfuric acid as a ‘diprotic’ or ‘dibasic’ acid and
carbonate as a ‘diacid’ base); strong and weak
acids and bases; amphoteric compounds;
dissociation constants Ka and Kb and their units .
Ionic product of water, Kw; the pK convention;
pH of aqueous solutions; hydrolysis of salts.
Students should be able to:
use pH to determine the [H3O+] and [OH-] of solutions
and vice versa;
distinguish between strong and weak acids and bases in
terms of the values of the constants concerned;
relate Kw to Ka and Kb;
calculate concentrations of species using dissociation
constants for weak acids and bases;
recognize that dibasic acids dissociate in a stepwise
manner with an associated Ka value for each step;
Indicators,
typified by phenolphthalein and
methyl orange.
describe use of pH meter and indicators to measure pH;
decide on use of appropriate indicators for titrations and
recall the pH range of phenolphthalein and methyl
orange;
use Ka values of indicators to justify their use in acid-base
titrations.
pH curves
Note: pH curves are restricted to monobasic acid –
monoacid base reactions.
Buffer solutions
Students are expected to be able to:
explain a buffer system and how it works to keep the pH
fairly stable;
explain applications of buffers (e.g. in blood, in seawater;
in industry);
perform calculations involving pH of buffer solutions.
8.2
Role of solvent in equilibria involving ionisation
of molecular solutes. Degree of ionisation.
Students should be able to:
Qualititative explanation of conductivity of
solutions of strong and weak electrolytes.
calculate the degree of ionization of weak electrolytes
which are acidic or basic;
recognize the connection between conductivity of a
solution and the presence of dissolved ions;
explain the special case represented by the conductivity
of the proton and the hydroxide ion in aqueous systems;
explain how to carry out and interpret the results of a
conductimetric titration (not involving weak acids with
weak bases).
8.3
Heterogeneous ionic equilibria: solubility
product Ksp, its units and relation to molar
solubility.
Students should be able to:
write solubility product expressions for sparingly soluble
salts;
solve mathematical equations involving Ksp and molar
solubility of ionic sparingly soluble solutes of the type
1:1 (e.g. AgCl or BaSO4);
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AM Syllabus (2015): Chemistry
Factors influencing
equilibrium.
the
solubility
product
describe the effect on the solubility product equilibrium
by addition (common ion effect) or removal of an ion
involved in the equilibrium by acid-base or complex ion
formation.
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AM Syllabus (2015): Chemistry
9
9.1
Redox Equilibria
Redox reactions: balancing of equations using
half equations.
Disproportionation reaction.
Oxidation number (state).
Students should be able to:
distinguish between oxidation number (state) and the
actual charges on atoms;
explain the connection between oxidation number and
bonding, e.g. for carbon in CO, CO2 and hydrocarbons;
oxygen in H2O and H2O2 sulfur in S2O32- and S4O62- ions;
state common oxidising and reducing agents as are used
in the laboratory and in the chemical industry.
9.2
Electrodes, galvanic cells and standard electrode
potentials. Standard hydrogen electrode as a
reference electrode.
Students are expected to be able to describe the hydrogen
electrode as a typical standard electrode.
9.3
Cell diagrams
Students should be able to:
draw a cell diagram for a galvanic couple, for example
for the copper-zinc cell, as follows:
Cu(s)│Cu2+(aq) ║ Zn2+(aq) │Zn(s) and determine for
such a cell the polarity and the emf;
The equation ∆G = -z EoF
describe the connection between the free energy change
of a reaction its Eo value and also employ electrode
potentials to determine whether redox reactions are
energetically possible.
Application of electrode potentials to the predict
redox change. Electrochemical series.
Students should be able to:
9.4
use the Nernst Equation to predict redox change under
non-standard conditions (although the equation will be
provided where its use is required);
recognize that a pH meter makes use of the Nernst
equation to measure the hydrogen ion concentration in a
solution (the theory behind the glass electrode will not be
tested);
describe theory of fuel cells using the hydrogen – air and
methanol – air cells as examples (operational details and
construction of fuel cells will not be tested).
Fuel cells
Corrosion as an electrochemical
sacrificial protection.
process:
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AM Syllabus (2015): Chemistry
10 Reaction Kinetics
10.1
The experimental investigation of reaction rate
for simple reactions.
Students should be able to describe suitable methods
(including their relative merits and demerits) for the
determination of reaction rates and be able to interpret
the data obtained.
The methods that shall be tested are:
gas volume changes by gas syringe;
colorimetry;
conductance measurements;
changes of pH for (slow) reactions involving
consumption or formation of hydroxonium ions;
titrimetry.
10.2
Order and rate coefficients and their
measurements, including the initial rate method;
rate equation.
Students should be able to:
define order of reaction and rate constant;
state the units of rate constants;
derive the rate equation for a reaction from data provided;
present and interpret kinetic data from graphical
representations or extract information from graphical
representations of kinetic data.
Integrated rate equations and fractional and pseudo orders
will not be tested.
Students should be able to:
Half-life for first order reactions.
state that the half life of a first order reaction is constant;
determine half life of a first order reaction from given
data and draw the decay curve from the half life.
10.3
10.4
Effects of pressure, concentration, surface area,
temperature and catalysts on reaction rate.
Activation energy of reactions.
MaxwellBoltzmann distribution of energies and the
collision theory of reactions.
Students should recognize the connection between the
reaction rate, temperature and/or activation energy as
represented by the Arrhenius equation.
Concept of rate-determining step in a multistep
reaction and its relevance to the mechanism of
the reaction. The molecularity of a reaction.
Students should be able to:
The equation will be provided when required and
numerical calculations will not be set.
distinguish between order and molecularity of a reaction;
describe examples of mechanisms with a slow step
followed by a number of fast steps.
10.5
Third order reactions and notion of the
improbability of three-particle collisions.
Mechanisms with a fast reversible first step followed by a
slow step will not be examined.
Photochemical
mechanisms.
Students should be able to describe initiation, propagation
and termination steps of free radical mechanisms typified
by the chlorination of methane, the polymerisation of
ethene and the formation of ozone in the stratosphere (an
important reaction in view of the role of ozone in filtering
ultraviolet radiation).
reactions
and
free
radical
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AM Syllabus (2015): Chemistry
10.6
Homogeneous and heterogeneous catalysis;
examples of the use of catalysts in industrial
processes; catalytic converters; autocatalysis.
Students should be able to:
distinguish between homogeneous and heterogeneous
catalysts;
describe the role of catalysts in speeding up a reaction by
providing an alternative pathway with a lower activation
energy;
describe and explain their lack of influence on the
equilibrium composition of a reversible reaction;
draw enthalpy level diagrams (energy profiles) for
reactions with and without catalysts;
describe the role of the catalytic converter in changing
unburnt hydrocarbons, CO and nitrogen oxides in car
exhaust into less harmful products (details of different
type of converters will not be tested);
describe the phenomenon of catalyst poisoning.
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AM Syllabus (2015): Chemistry
11 Chemical Periodicity
Periodic classification in terms of electronic structure.
Periodic relationships amongst the elements Li to Ar.
Students should be able to:
describe how properties of elements exhibit periodic
behaviour as illustrated by their trends in melting and
boiling points, electrical conductivity and ionisation
energy; and by their reaction with oxygen, chlorine and
water;
write formulae of the oxides and describe their physical
state and relate it to their structure and bonding and state
the acid-base character of oxides/hydroxides of metals
and oxides of non-metals for the following Na2O, MgO,
Al2O3, NaOH, Mg(OH)2, Al(OH)3, SiO2, P4O10, SO2,
SO3, Cl2O;
write formulae of chlorides (Li to N and Na to P),
describe their physical state and relate it to their structure
and bonding and state the reactions of chlorides with
water;
write formulae of the simple hydrides (Li to F and Na to
Cl but excluding boron and aluminium hydrides),
describe their physical state and relate it to their structure
and bonding.
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AM Syllabus (2015): Chemistry
12 Inorganic chemistry
Comparative study of the s-block elements (i) lithium,
sodium and potassium; (ii) beryllium, magnesium,
calcium, strontium and barium.
12.1
Fixed
oxidation
states
and
configuration in terms of orbitals.
electronic
Physical properties of elements; flame tests;
reaction with water and with oxygen.
Diagonal relationships: Li and Mg; Be and Al
(see also section 12.2 on Al).
Oxides, peroxides, superoxides, hydroxides,
nitrides, carbonates, hydrogencarbonates,
chlorides, nitrates(V), nitrates(III), sulfates(VI)
and sulfates(IV).
Students should be able to explain the reasons for
diagonal relations and give examples to illustrate the
relation.
Students should be able to:
draw structures of these ions, including any covalent
bonds and any delocalization features;
draw the crystal lattices of NaCl and CsCl and to state the
coordination number of the metal ion in these chlorides;
Trends in thermal stability of nitrates(V) and
carbonates and solubility of hydroxides and
sulfates(VI).
12.2
Aspects of the chemistry of aluminium: hydrated
aluminium ion, oxide and hydroxide; chlorides,
alums, nitrates.
describe a method of preparation of the following
compounds: oxides of Li and group II metals; peroxides
of Na and Ba; KO2; Mg(OH)2 by precipitation reactions;
describe the effect of heat on nitrates(V),
hydrogencarbonates and carbonates;
explain trends in solubility on the basis of hydration and
lattice enthalpies.
Students should be able to explain:
the amphoteric nature of the metal and its oxide and
hydroxide;
passivation of the metal;
structure and acidity of the hydrated ion;
a method of preparation of anhydrous AlCl3 (by direct
union of the elements) and hydrated aluminium chloride;
the structure of the dimeric form of gaseous Al2Cl6 and
know that it exists in equilibrium with the monomeric
form;
alums as double salts of general formula
MIMIII(SO4)2.12H2O where MIII could be an ion other
than aluminium (preparation of alums is not required).
12.3
Chemistry of carbon, silicon, tin and lead.
Students should be able to:
Allotropy of the elements
describe the structure of the allotropes of carbon
(graphite, diamond and fullerenes typified by C60) and the
properties of diamond and graphite and relation to
structure and bonding;
Main oxidation states and inert pair effect
explain the variation in ionization energies;
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AM Syllabus (2015): Chemistry
predict the relative stability of the two main oxidation
states on descending the group;
explain the redox character of simple compounds of
Sn(II) and Pb(IV);
Principal oxides
predict the nature of the bonding in a given oxide and
describe the acid/base character of oxides including the
amphotericity of tin(II) and lead(II);
Hydrides
state how CH4 can be produced in the lab (by
decarboxylation reaction or by hydrolysis of Al4C3) and
how SiH4 can be made from Mg2Si and acid. Predict
relative stability towards heat;
Halides: preparation and hydrolytic behaviour;
formation of complex ions.
describe a method of preparation of the halides
(excluding fluorides) of the elements (both in the IV and
II states);
know how tetra- and dichlorides (restricted to Sn and Pb)
react with water and recognize why CCl4 does not
hydrolyse;
describe the formation of PbCl42-, SnCl62-;
Toxicity of lead
explain the use of lead tetraethyl in petrol engines and
why it has been phased out.
12.4
Chemistry of fluorine, chlorine, bromine and
iodine.
Students should be able to:
Laboratory preparation of chlorine, bromine and
iodine and simple tests for these elements.
describe one method of preparation including the
practical details; recall the bleaching property of chlorine
and bromine, the colour of solutions of bromine and
iodine in organic solvents; the starch test for iodine;
Trends in physical and chemical properties for
the series X2 and HX (X = F, Cl, Br and I).
describe and explain the trends in boiling points and
bond dissociation enthalpies;
Trends in the ionisation energies and electron
affinities of the halogens.
-
Reactions between ionic halides X (X = Cl, Br,
I) and silver ions, phosphoric(V) and
sulfuric(VI) acids.
explain why the solubility of AgX (X = Cl, Br, I) is
influenced in the presence of ammonia and thiosulfate
and recognize the formation of the corresponding silver(I)
complexes;
explain the reaction of the halide ions with concentrated
sulfuric(VI) acid in relation to the reducing properties of
X- and HX;
Acidity of HX (X = F, Cl, Br, I) and anomalous
behaviour of HF.
explain the trend in acid strength for HX and describe and
explain the hydrogen bonding in HF and in the ion HF2-;
Oxidation states of the halogens Cl, Br and I in
their compounds and the relative oxidising
strength of halogens X2 (X = Cl, Br, I) . Oxo
anions as oxidising agents.
write ionic equations describing the disproportionation
reactions of chlorine and chlorate(I) (OCl-);
describe the structures of chlorate(V) and chlorate(VII)
ions;
describe the effect of heat on KClO3;
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AM Syllabus (2015): Chemistry
describe the oxidation of iodide by chlorate(V);
12.5
Iodine in titrimetry; triiodide ion.
The chemistry of the transition elements.
describe the reaction of iodine with thiosulfate ion and
the use of this reaction in redox titrimetry.
Students should be able to:
Difference between transition and d-block
elements.
identify as a ‘transition element’ that which forms at least
one species in which the d-orbital is incompletely filled;
A general overview of the transition metals
emphasising features they have in common.
give the electronic configurations of the metals Ti
through to Cu (given their atomic number);
account for their metallic character, variable oxidation
states explained on the basis of the electron configuration
of these elements, catalytic roles, formation of coloured
compounds and formation of complex ions - explanation
of colour in terms of electronic transitions between dorbitals of different energy will not be examined;
12.6
Complex ions
describe bonding in complex ions in terms of the
electrostatic model or else in terms of dative covalency
between ligand and central metal ion;
Shapes of di-, tetra-, and hexa-coordinated
systems; isomerism in complex ions.
describe and illustrate structural and stereoisomerism in
octahedral complexes;
Stability of complex ions and ligand exchange.
write equilibrium reactions related to stepwise formation
constants and predict ligand exchange from values of
formation constants;
Stability of chelating ligands.
know the bidentate ligands ethylenediamine, ethanedioate
and glycinate ion and the hexadentate ligand EDTA4- and
describe how these ligands easily displace monodentate
ligands to form more stable chelate complexes;
Nomenclature of complex ions.
name complex ions (using as a guide the publication
Chemical Nomenclature for Use in Matriculation
Examinations).
More detailed chemistry of the following dblock elements: chromium, manganese, iron and
copper.
Chromium:
Students should be able to :
Reaction of metal with acids, including
passivation with HNO3, with alkalis to form
chromate(III).
write relevant chemical equations to describe these
changes and explain passivation;
Chromium(II) state and its strong reducing
character.
describe, using a balanced chemical equation, the
formation of Cr(II) from Cr(VI) by reduction with Zn/H+;
Chromium(III) compounds
describe how to prepare from readily available precursors
(e.g. simple salts) the oxide and hydroxide of Cr(III) and
chromates(III);
describe the conversion of Cr(III) to chromate(VI);
write the possible isomers from the hydrates of CrCl3
(method of preparation of isomers will not be tested);
recognize that Cr(III) can form alums;
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AM Syllabus (2015): Chemistry
Chromium(VI) compounds
describe and illustrate the structures of the chromate(VI)
and dichromate(VI) ions;
describe the chromate(VI)-dichromate(VI) equilibrium
and the effect of pH change on this equilibrium;
describe the use of chromate(VI) in qualitative analysis as
a precipitant and dichromate(VI) in titrimetry and as an
oxidant in organic chemistry;
know how ammonium dichromate(VI) decomposes on
heating;
know how CrO3 can be made and how it can be converted
into chromic(VI) acid;
know about the toxicity of Cr(VI) compounds.
Manganese:
Students should be able to:
Reactivity of metal exhibited by its reaction with
dilute acids to form Mn2+ and hydrogen.
describe the oxide and hydroxide of Mn(II) and explain
how Mn(OH)2 oxidises easily to Mn(III) and Mn(IV);
Manganese(II) compounds
describe how manganese(II) can be oxidised
manganese(VII) using bismuthate and persulfate ion;
Manganese(IV) and manganese(VI) compounds
recall that MnO2 is a dark brown solid that can be
prepared from Mn(II) in the lab and that it is a strong
oxidising agent capable of oxidising HCl to chlorine,
ethanedioate to carbon dioxide and alkylbenzenes to the
corresponding aldehyde;
to
describe how MnO2 can be converted into manganate(VI)
and how this species disproportionate;
recall the tetrahedral structure of MnO4–;
Manganese(VII) compounds
explain the use of KMnO4 as a strong oxidising agent in
titrimetry and in organic chemistry.
Iron:
Students should be able to:
Reaction of metal with oxidising and nonoxidising acids.
describe Fe3O4, (which occurs naturally as magnetic iron
ore) as an example of a mixed oxide containing iron(II)
and iron(III) ions;
Reaction of iron with steam to produce hydrogen
and Fe3O4 (iron(II) diiron(III) oxide)
describe the rusting of iron as an electrochemical process;
explain how rusting can be prevented including the use
of sacrificial anodizing;
Reaction of iron with non-metals including C to
form steels; rusting of iron
describe methods of formation of the oxide (a very
unstable, easily oxidisable substance), hydroxide, iron(II)
sulfate-7-water and the double salt ammonium iron(II)
sulfate-6-water;
Iron(II) compounds;
compounds.
iron(III)
describe methods of oxidizing iron(II) compounds;
describe methods of formation of the oxide and of
anhydrous FeCl3;
iron(II)
describe the reaction of iron(III) with alkali, including
aqueous ammonia;
oxidation
Iron(III) compounds; reduction
compounds.
to
to
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AM Syllabus (2015): Chemistry
describe methods of reducing iron(III) compounds;
give reasons for the non-existence of iron(III) iodide and
iron(III) carbonate;
explain the acidity of hexaaquairon(III) ion and compare
with that of the corresponding iron(II) species;
Complex ions of iron
describe the formation of [Fe(CN)6]4- by reaction of
iron(II) and cyanide and explain how it can be converted
into [Fe(CN)6]3- by oxidation with chlorine;
explain the use of hexacyanoferrates in qualitative
inorganic analysis;
state how [Fe(SCN)(H2O)5]2+ is involved in qualitative
analysis for Fe(III);
describe the ion [Fe(NO)(H 2O)5]2+ as an Fe(II)-NO
complex and explain how it forms during the brown ring
test for nitrate(V).
Copper:
Students should be able to:
Reaction with oxidising acids
write chemical equations for reactions of Cu with
nitric(V) acid of various concentrations and with
concentrated sulfuric(VI) acid;
Copper(I) compounds
describe how Cu2O is formed in Fehling’s test for
aldehydes and how this oxide can be converted into CuCl
and CuI;
describe the instability of Cu+(aq) with respect to
disproportionation and to explain how ligands other than
water stabilize this ion;
Copper(II) compounds
describe methods of preparation of the oxide, hydroxide,
sulfate(VI), nitrate(V) and basic carbonates of Cu(II);
Complex ions of copper(I) and copper(II)
explain the formation of complex ions of Cu(I) and Cu(II)
with ammonia and chloride ions, the formation of the
glycinate complex of Cu(II);
describe the shape of Cu(II) complexes as either square
planar or distorted octahedral species; tetrahedral CuCl42–
;
explain the interconversion of blue aqua copper(II)
species into yellow chlorocuprate(II).
12.7
A more detailed study of the inorganic chemistry
of the following non-metals.
Hydrogen
Students should be able to:
Manufacture of hydrogen
describe the manufacture of the element from natural gas
by steam reformation and by electrolysis of water;
Laboratory preparation
explain methods of preparation of hydrogen on a small
scale using reactions involving metal + acid or Al +
alkali;
Properties as reducing agent
describe reactions of hydrogen acting as a reducing agent
with non-metals (oxygen, nitrogen and sulphur) and metal
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AM Syllabus (2015): Chemistry
oxides as well as with organic molecules, e.g. unsaturated
hydrocarbons;
know about the old concept of ‘nascent hydrogen’ and its
modern explanation in terms of reduction by electrons
released in the metal-acid reaction;
Ionic, covalent and interstitial hydrides
Complex hydrides: LiAlH4 and NaBH4
describe the bonding in salt-like hydrides (containing Hion) and covalent hydrides and explain the connection
between interstitial hydrides and catalysis in
hydrogenation reactions;
explain the structure and bonding in LiAlH4 (lithium
tetrahydridoaluminate) and how it can be made in the
laboratory;
use of LiAlH4 as a reducing agent in organic chemistry
for the >C=O and the -COOH groups but which does not
reduce the C=C double bond;
Water and its physical properties; solvent action
account for the physical properties of water in terms of
hydrogen bonding and for its solvent properties in terms
of its polar structure
Deuterium oxide and deuteroderivatives (eg:
DCl, ND3, C2D2).
Nitrogen
describe how simple deuteroderivatives can be
synthesised using D2O as a source of deuterium (the
method of formation of deuterium oxide will not be
tested).
Students should be able to:
Structure and bonding of N2
explain the lack of reactivity of nitrogen as a consequence
of the strength of the triple bond;
Ammonia and its reactions as base, ligand and
reducing agent; ammonium salts
describe the properties of ammonia as a weak LowryBrønsted base, a good monodentate ligand (with Cu2+,
Zn2+, Ni2+ in qualitative analysis) and a reducing agent;
explain the thermal instability of ammonium nitrate(III)
and ammonium nitrate(V) as resulting from the redox
reaction between the reducing cation and the oxidising
anion;
Oxides of nitrogen N2O, NO, NO2, N2O4, N2O5
describe a laboratory preparation of these oxides of
nitrogen;
draw electronic structures to show the bonding in these
oxides and to recognise the presence of odd electron
systems;
draw molecular shapes of these oxides;
describe the disproportionation of NO2 when dissolved in
water or alkali;
Nitric(III) (nitrous) acid and salts
describe the laboratory preparation of nitric(III) acid and
nitrates(III) of s-block elements;
Nitric(V) (nitric) acid and salts
describe the laboratory preparation of nitric(V) acid by
reaction of an involatile acid on nitrates(V);
describe reactions of nitric(V) acid acting as an acid, an
oxidising agent and a nitrating agent in organic
chemistry;
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AM Syllabus (2015): Chemistry
describe the reduction of nitrate(III) and nitrate(V) to
ammonia using ‘nascent hydrogen’ type systems and the
brown ring test for nitrate(V).
Oxygen
Students should be able to:
Allotropes of oxygen: O2 and O3 (trioxygen or
ozone)
state that oxygen is a main component of the atmosphere;
test for oxygen in the lab;
explain how trioxygen forms in the stratosphere from
photochemical reactions and its action as a screen of
ultraviolet radiation;
state that trioxygen forms in the lower atmosphere as an
air pollutant (chemistry of formation will not be tested);
Laboratory preparation of oxygen
describe the preparation of oxygen gas from oxides and
peroxides;
Reaction of oxygen with metals and non-metals
describe the formation of compounds of oxygen by direct
union between oxygen and various elements and to
classify oxides as acidic, basic, amphoteric or neutral;
distinguish between normal and mixed oxides (typified
by Fe3O4 and Pb3O4) in terms of the species present in
these compounds;
describe a laboratory preparation of hydrogen peroxide
from barium peroxide;
Hydrogen peroxide and peroxides
describe the main properties of hydrogen peroxide
namely, as oxidising agent, as a reducing agent with
oxidants stronger than itself (MnO4– and Cr2O72–), and as
a thermally unstable liquid;
recognize the relatively weak nature of the peroxide bond
and the ease of its homolysis to form free radicals;
describe the structure and bonding of the peroxide ion in
sodium and barium peroxide (as typical compounds) and
its reaction with water, acid and with carbon dioxide;
Superoxide ion
describe the structure and bonding of the superoxide ion
in KO2 and its reaction with water.
Sulfur
Students should be able to:
Allotropes of sulfur
monoclinic sulfur
(S8):
rhombic
and
Hydrogen sulfide
recognize the effect of temperature on the stability of the
two allotropes;
recognize the odour and toxicity of hydrogen sulfide;
describe how H2S can be prepared from sulfides and
explain its properties as a weak dibasic acid, a reducing
agent and as a precipitant for insoluble sulphides;
Sulfur dioxide and sulfuric(IV) (sulfurous) acid,
sulfates(IV) (sulfites)
describe how sulfur dioxide, sulfuric(IV) and sulfates(IV)
can be made and interconverted in the laboratory and
explain their reducing character;
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AM Syllabus (2015): Chemistry
Sulfur trioxide and sulfuric(VI) (sulfuric) acid
describe how sulfur trioxide, sulfuric(VI) and sulfates(VI)
can be made and interconverted in the laboratory;
describe reactions of concentrated sulfuric(VI) acid as an
involatile proton donor capable of displacing more
volatile acids from their salts, oxidising agent, strong
dehydrating agent and sulfonating agent in organic
chemistry;
Thiosulfate
describe the formation of thiosulfate ion from sulfate(IV)
and its reaction with acids, iodine and chlorine;
Structure of oxoanions of sulfur
draw the electronic structures and shapes of the ions:
sulfite, sulfate and thiosulfate and recognize the presence
of electronic delocalization in these structures;
recognize the fact that the two sulfur atoms in thiosulfate
are present in different structural environments and the
implication of this on the oxidation state;
Sulfur oxides as environmental pollutants
explain why sulfur oxides are often present in the
atmosphere as contaminants, their origin in fossil fuels
and their role in the formation of acid rain and its
environmental impact.
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AM Syllabus (2015): Chemistry
13.1
Organic chemistry: general principles
13.1
Functional groups and homologous series.
molecular, empirical and structural formulae
including graphical representations
Students should be able to:
distinguish between molecular and empirical formulae
and write possible structural formulae from molecular
formulae;
identify functional groups in the following classes:
alkenes, alkynes, arenes, , alcohols, ethers, phenols,
halogenoalkanes and halogenoarenes, aldehydes,
ketones, carboxylic acids and derivatives, amines and
diazonium compounds;
recognize as structural formulae systems such as
CH3CH2CH3 or CH3CHO as well as
graphical
(displayed) formulae where individual bonds are shown
explicitly;
use wedge and dashed/hatched line convention to
represent three dimensional structure;
Nomenclature of organic compounds
give systematic names of organic compounds from their
structure and vice versa (refer to publication Chemical
Nomenclature for Use in Matriculation Examinations).
13.2
Purification of compounds
Students should be able to:
describe the following techniques used in the purification
of
organic
compounds:
solvent
extraction,
recrystallisation, drying (for gases, liquids and solids),
simple, fractional and steam distillation, sublimation;
column, and thin layer chromatography (one
dimensional);
Preparation of derivatives and their use for
characterisation and purification
describe the process of derivatisation for characterization
purposes (in aldehydes, ketones and alcohols) and for
purification
purposes
(hydrogensulfite
addition
compounds for aldehydes);
Determination of melting points as a test for
purity
recognize that presence of impurities cause a melting
point to deviate from the standard value;
Mixed melting point technique
describe the mixed melting point technique as a method
of identification of compounds.
13.3 The determination of empirical, molecular and
structural formulae from analytical information.
Experimental techniques are not required but students
should be able to use percentage composition data to
deduce the empirical formula and, through relative
molecular mass, obtain the molecular formula.
13.4 Structure deduced from chemical and physical
properties and from information derived from
instrumental techniques, namely, mass spectrometry and
infra red spectroscopy.
Students should be able to:
deduce the presence of specific functional groups from
results of reactions with reagents;
interpret infra red information (including spectra) for the
identification of simple functional groups using given
tables of frequencies/wave numbers (details of
instrumentation are not required).
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AM Syllabus (2015): Chemistry
13.5
Catenation in carbon compounds and spatial
distribution of the bonds in simple carbon compounds.
Students should be able to:
explain why the element carbon is unique in exhibiting
catenation;
explain the shape of the bonds and bond angles in alkanes
(tetrahedral), alkenes and carbonyl groups (planar) and
alkynes (linear) in terms of hybridized sp3, sp2 and sp
orbitals.
13.6
Isomerism in organic compounds
Students should be able to:
describe structural and stereoisomerism namely: cis-trans
(geometrical) restricted to C=C and C=N systems;
optical isomerism (restricted to enantiomerism) and
conformational isomerism (restricted to cyclohexane);
describe the property of optical activity in terms of
dextrorotatory (+) and laevorotatory (–) compounds and
connect it to chiral molecules being those which do not
possess a plane of symmetry (including systems which
possess two chiral carbons but still having a plane of
symmetry), use of the descriptors D and L and R and S
are not required;
explain the term racemate/racemic mixture of optical
isomers and distinguish between internal and external
compensation (resolution of racemic mixture will not be
tested);
describe the phenomenon of keto-enol tautomerism and
be able to predict its occurrence in structures that may be
novel to them;
explain how size, shape and polarity of molecular
isomers can influence physical properties including
comparison of volatility of linear and branched isomers;
describe the effect of intramolecular versus
intermolecular hydrogen bonding on boiling point as
typified by the nitrophenols;
13.7
Delocalisation of electrons in organic molecules
Students should be able to:
describe the benzene ring in terms of sigma () and pi ()
bonds and to recognize the significance of delocalization
in this and similar molecules;
predict the presence of delocalisation in conjugated
systems and know that delocalization lowers the energy
of the system and is thus a favourable structural feature;
use canonical formulae, including Kekule structures for
aromatic molecules, for the description of delocalization
in organic molecules.
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AM Syllabus (2015): Chemistry
14
14.1
Chemistry of Aliphatic and Aromatic Compounds
Petroleum: hydrocarbons from petroleum
Students should be able:
to recall that alkanes, alkenes and aromatic hydrocarbons
can be obtained
from
petroleum by fractional
distillation;
to describe the principle of the processes of cracking and
reforming (no technical details of the processes are
expected);
Use of hydrocarbons as fuels
to explain how the combustion reactions of alkanes and
other hydrocarbons lead to their use as fuels in industry,
in the home and in transport;
Environmental considerations
to explain how the combustion of hydrocarbons leads to
the formation of carbon dioxide, an important greenhouse
gas;
14.2
Alkanes, alkenes and alkynes
to explain the formation of carbon monoxide, oxides of
nitrogen and unburnt hydrocarbons arising from the
internal combustion engine and the use of catalysts to
minimise their effect on pollution.
Students should:
Laboratory preparation of alkanes
know that alkanes can be obtained from a number of
organic compounds as indicated elsewhere in this
syllabus;
Alkanes and cycloalkanes: reactivity and
reactions
be aware of the general unreactivity of alkanes, including
towards polar reagents;
be able to describe by appropriate chemical equations
free radical substitution by chlorine and bromine;
Concept of ring strain
be able to describe how ring strain in molecules, such as
cyclopropane, leads to enhanced reactivity relative to
corresponding non-cyclic alkanes.
Alkenes: alkenes as examples of unsaturated
hydrocarbons
Students should be able to recognise:
alkenes as unsaturated hydrocarbons containing a double
bond made up of a σ and π bond;
the double bond as an electron rich centre susceptible to
electrophilic attack.
Preparation of alkenes
Students should be able to describe the preparation of
alkenes from alcohols and haloalkanes through
elimination reactions.
Addition reactions of alkenes
Students should be able to:
describe the reaction of alkenes with halogen, hydrogen,
halogen halides, and water in presence of sulfuric(VI)
acid;
Markovnikov’s rule
identify the major product in the reaction of halogen
halides with an unsymmetrical alkene and explain its
formation in terms of relative carbocation stability;
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AM Syllabus (2015): Chemistry
Ozonolysis as the addition reaction of ozone to a
double bond followed by hydrolysis in the
presence of zinc/ethanoic acid
draw the structure of the unstable ozonide;
describe the reaction as producing carbonyl compounds
as products of hydrolysis of the ozonide;
describe the use of the reaction for the location of double
bond;
Oxidation of the double bond
describe the reaction with cold,
manganate(VII) ions to form the diol;
Addition polymerisation
describe the following characteristics of addition
polymerisation such as that involved in formation of
poly(ethene) and PVC: deduce the repeat unit of a
polymer obtained from a given monomer;
dilute,
alkaline
identify the monomer present in a given section of a
polymer molecule;
Hydrocarbon poly(alkene)s as waste products
recognise the use of poly(alkene)s waste in recycling and
as a fuel and also the nonbiodegradability of these
products;
Alkynes
Preparation of alkynes
recognise alkynes as unsaturated hydrocarbons containing
a triple bond made up of a σ and two π bonds;
describe
the
preparation
dehydrohalogenation reactions;
of
alkynes
by
describe the preparation of ethyne from CaC2;
14.3
Addition reactions of alkynes
state the reactions with hydrogen, X2 and HX to form
partial or fully saturated products and with water to form
carbonyl compounds via tautomerisation of the enol
(students should recognize the importance of
chloroethene as a feedstock for PVC manufacture);
Combustion
describe the exothermic reaction of ethyne with oxygen
and its use in oxy-acetylene flame;
Acidity of terminal alkynes
describe the reaction of ethyne and propyne with
ammoniacal silver nitrate and copper(I) chloride.
Arenes: benzene and alkylbenzenes
Students should:
Electrophilic substitution
be able to describe the structure and bonding in benzene
and similar aromatic systems;
be able to describe how to introduce the nitro, sulfonate,
alkyl, alkanoyl and the halogen group on the aromatic
ring;
be familiar with the concept of activating and
deactivating groups and their directing influence on
further substitution of the aromatic ring. Such groups are
to be exemplified: -CH3, -OH, -NH2, -X, -CHO, -COOH;
Addition reactions
be able to describe the addition of hydrogen and chlorine
to benzene under appropriate conditions;
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AM Syllabus (2015): Chemistry
Side chain reactions of methylbenzene:
(i)
halogenation
describe the free radical substitution of the methyl side
chain with chlorine or bromine;
(ii)
be able to convert the –CH3 side chain to –CHO and
–COOH using chromium(VI) dichloride dioxide and
manganate(VII) respectively;
oxidation
recognise the fact that any alkyl group is oxidised to –
COOH upon oxidation with manganate(VII).
14.4
Alcohols, ethers, phenols
Students should be able to:
Classification of alcohols
identify a given alcohol as primary, secondary or tertiary
and distinguish between classes in terms of oxidation and
triiodomethane reaction;
know about the Lucas test and recognize that it can easily
distinguish tertiary alcohols from the other two classes,
(the test has dubious validity in distinguishing secondary
from primary alcohols and for this reason it will not be
employed in the practical examination for this purpose);
Physical properties of alcohols
explain the relatively high boiling point and miscibility
with water of lower members in terms of hydrogen
bonding in comparison to ethers;
Preparation of alcohols
describe how to prepare alcohols from halogenoalkanes,
alkenes, by reduction of aldehydes, ketones and
carboxylic acids, from Grignard reagents;
Reactions of alcohols
describe the formation of alkoxides, halogenoalkanes,
esters, alkenes and ethers from alcohols;
Aromatic alcohols
distinguish
aromatic
alcohols,
phenylmethanol, from phenols;
Polyhydric alcohols
recall ethane-1,2-diol and propane-1,2,3-triol as typical
polyhydric alcohols;
Industrial preparation of ethanol
describe the formation of ethanol by the fermentation of
sugars and by steam hydration of ethane;
Ethers
recall the preparation of ethers from alkoxides and
halogenoalkanes;
typified
by
describe the effect of HI on ethers;
recognise the use of ethoxyethane as a solvent and in
extraction procedures despite its inflammable nature;
Phenols
Laboratory preparation:
describe the preparation of phenol from sulfonic acids
and diazonium salts;
Acidity of phenols:
recognise the weak acid nature of phenol by its failure to
liberate CO2 with sodium carbonate and sodium
hydrogencarbonate (contrast with carboxylic acids);
describe the formation of esters and ethers from phenols;
Reactions of phenol, including substitution
reactions
describe substitution reactions in the ring to include
tribromination, nitration;
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AM Syllabus (2015): Chemistry
describe the reduction to benzene and the coupling
reactions with diazonium ions to form azo dyes;
recall that phenol and related compounds give a
characteristic violet colouration with neutral iron(III)
chloride which serves as a diagnostic test for phenols.
14.5
Halogenoalkanes and halogenoarenes
Students should be able to:
Preparation
describe the preparation of
halogenoalkanes from alcohols and alkenes;
Reactions
describe the conversion of halogenoalkanes into
alcohols, ethers, amines, nitriles, esters, alkanes, alkenes,
alkynes, Grignard reagents and alkylarenes;
Effect of structure on chemical properties
explain the reactivity of halogenoalkanes in terms of their
structure (primary, secondary, tertiary halogenoalkanes)
and the halogen atom (Cl, Br, I);
Halogenoarenes
14.6
Preparation
describe the preparation by direct halogenation (where
appropriate) and via diazonium compounds;
Effect of structure on chemical properties
explain the unreactivity of the substituent halogen atom
with respect to replacement reactions in terms of
delocalization of electrons;
(Chloromethyl)benzene
explain why (chloromethyl)benzene is more susceptible
to replacement reactions than nuclear halogenoarenes;
Reaction with magnesium
recall the formation of Grignard reagents;
Uses of organohalogen compounds other than in
synthetic organic chemistry
recall the use as solvents; explain the impact on the ozone
layer caused by the use of CFCs and other halogeno
compounds.
Aldehydes and ketones
Students should be able to:
Preparation
describe the preparation of aldehydes and ketones from
corresponding alcohols;
describe the preparation of aromatic aldehydes by
oxidation of methylarenes using chromium(VI) dichloride
dioxide and the preparation of aromatic ketones by
Friedel-Crafts reaction;
Addition reactions of the carbonyl group
describe the addition reactions with H- (as in the
hydridometallates); ROH; HCN; and NaHSO3 (including
the use of this reagent in the purification of aldehydes);
Condensation reactions
describe the reactions with hydroxylamine; hydrazine,
phenylhydrazine and 2,4-dinitrophenylhydrazine;
explain the use of 2,4-dinitrophenyl-hydrazones for
characterisation and identification purposes;
describe the formation of condensation polymers from
methanol;
recall the reactions with halogen and with PCl5; the
haloform reaction; the aldol reaction; the Cannizzaro
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AM Syllabus (2015): Chemistry
reaction; the reduction using zinc and acid and with
hydrogen;
14.7
Other reactions
distinguish aldehydes and ketones on the basis of
oxidation employing weak oxidant systems typified by
Fehling reagent and Tollen’s reagent.
Carboxylic acid
Students should be able to:
Formation of carboxylic acids:
describe the formation of carboxylic acids from alcohols
and aldehydes by oxidation and from carboxylic acid
derivatives by hydrolysis;
formation from alkylbenzenes by oxidation of the side
chain;
Weak acidity
explain the weak protic acidity of carboxylic acids in
terms of their structure and relative stability of the
carboxylate anion;
Effect on acidity of α-substitution by halogen
atoms or alkyl groups
account for the increase or decrease in acidity of
halogen, or alkyl substituted carboxylic acids in terms of
inductive effect;
compare the acidity of ethanoic acid with that of ethanol
and phenol and account for differences in terms of
structure;
Salt formation and alkalinity of carboxylate
anion
describe methods of formation of carboxylate salts and
explain the alkalinity through hydrolysis reactions;
Soaps
recall that sodium or potassium carboxylates having long
alkyl chains can act as soaps;
Decarboxylation reaction
Reduction to primary alcohols
Exceptional behaviour of methanoic acid
Formation of the following derivatives of
carboxylic acids: esters, acid chlorides and acid
anhydrides (restricted to simple anhydrides), acid
amides and nitriles (but excluding isonitriles).
be familiar with the use of soda lime to bring about
decarboxylation of carboxylic acid salts and to represent
the equation involved in terms of reaction with “NaOH”;
describe the conversion of –COOH to –CH2OH using
LiAlH4;
recall the anomalous behaviour of methanoic acid and its
salts as typified by reaction with soda lime, dehydration,
reaction with PCl5 and oxidising agents;
recall that esters can be made using the reaction of acids
with alcohols; acid chlorides or anhydrides with alcohols
or phenols;
recall the formation of esters from reaction of
halogenoalkane with silver carboxylate;
recall the formation of acid chlorides from reaction of
carboxylic acids with phosphorus chlorides and SOCl2
(sulfur oxide dichloride) and acid anhydrides from the
acid chloride;
describe the formation of carboxylic acid amides,
–CONH2, by thermal decomposition of ammonium
carboxylates and reaction of esters, acid chlorides and
anhydrides with ammonia;
describe the reactions of acid chlorides and anhydrides
33
AM Syllabus (2015): Chemistry
with primary
and secondary amines
corresponding N-substituted amides;
to
form
explain the formation of nitriles by dehydration of amides
using phosphorus(V) oxide;
recall that nitriles can also be made from reaction of
halogenoalkanes and KCN in ethanol solution and
recognise this reaction as a step-up technique in organic
synthesis;
recall that aromatic nitriles can be formed from
diazonium salts;
Reactions of derivatives
recall the acid and alkaline hydrolysis of the various
derivatives;
explain the relative ease of hydrolysis of acyl chlorides
and halogenoalkanes ;
compare ammonolysis
reactions;
reactions
with
hydrolysis
explain the reactions of acid chlorides and anhydrides
with alcohols;
explain the Hofmann degradation of amide and appreciate
its importance in step down synthesis;
recall the reduction of acid amides and nitriles to amines;
recall the reduction of esters to alcohols (reduction of
acid chlorides and anhydrides will not be tested);
recall the reaction of amides with nitric(III) acid;
Halogenoacids
explain the formation of halogenoacids by reaction of the
acid and halogen;
explain how amino acids can be prepared from
halogenoacids (practical details of the isolation of the
amino acid will not be tested);
Polyesters and polyamides
explain the formation of polyesters and polyamides and
state their major commercial use;
explain the environmental advantage of the ease of
hydrolysis of both the ester and peptide link and contrast
this with the properties of poly(alkene)s;
Dicarboxylic acids:
recall the use of ethanedioic acid and its salts as standard
reducing agents in titrimetry;
explain the formation of the acid anhydride from
dicarboxylic acids that allow it on the basis of
stereochemistry.
14.8
Amines and diazonium compounds
Students should be able to:
Structure
identify a given amine as primary, secondary or tertiary;
identify quaternary ammonium compounds;
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AM Syllabus (2015): Chemistry
Formation of amines
prepare amines by Hofmann degradation, reduction of
simple and N-substituted amides, reduction of nitriles and
hydrolysis of N-substituted amides;
prepare aromatic amines by reduction of nitroarenes by
tin and concentrated hydrochloric acid;
Basicity of amines
explain the relative basic strength of ammonia, primary
and secondary aliphatic amines and phenylamine in terms
of their structure;
explain the difference in basic strength between amines
and acid amides;
Reactions of amines
give the formation of alkylammonium salts from
corresponding amines;
explain the alkylation reaction (Hofmann reaction) and
acylation of amines using acid chlorides and anhydrides;
recognise phenylamine as an activated ring and use its
reaction with aqueous bromine as an example of this
reactivity;
explain the reaction of nitric(III) acid with primary
aliphatic amines to produce nitrogen as a product of
decomposition of unstable aliphatic diazonium
compounds and recognize that this reaction is useful as a
test for primary aliphatic amines.
Diazonium salts
Students should be able to:
draw structures of diazonium salts and recognize their
ionic character;
Formation of diazonium compounds
14.9
describe the preparation of aryl diazonium salts by
reaction with dilute hydrochloric acid and sodium
nitrate(III) in ice cold conditions, account for the relative
stability of the the diazonium ion on the basis of
delocalization of charge;
Reactions of diazonium compounds:
(i) substitution reactions in which the diazo
group is lost
convert the diazonium salt to phenols, halogenoarenes,
arenes and nitriles;
(ii) reaction in which the diazo group is retained
give coupling reactions with phenols to form azo dyes.
Amino acids
Students should be able to:
recognize amino acids as typical of
difunctional
molecules which are important in biological systems;
Properties of amino acids
predict the reactions of amino acids as those pertaining to
the carboxyl and amino groups including esterification,
acylation, etc.;
know about the formation of dipolar ions (zwitterions) in
amino acids and the influence on physical properties
such as solubility in water and organic solvents and
melting points;
explain the existence of optical activity of amino acids;
35
AM Syllabus (2015): Chemistry
describe the effect of varying pH on the composition of
an aqueous solution of an amino acid: isoelectric point;
(technique of electrophoresis shall not be examined);
explain the reaction of amino acid salts as bidentate
ligands as typified by the copper(II) complexes;
describe a practical method of separation of a mixture of
amino acids using chromatography;
relate amino acids to peptides and proteins;
14.10
Preparation of amino acids
describe the synthesis of amino acids from carboxylic
acids or aldehydes via cyanohydrins.
Mechanistic aspects
Students should be able to:
use reaction mechanisms to rationalise the facts of
organic chemistry;
Ionic mechanisms
recall the following mechanistic concepts: nucleophile,
electrophile, free radical; inductive and mesomeric
effects;
distinguish between bimolecular and unimolecular
nucleophilic substitution reactions of halogenoalkanes;
describe electrophilic substitution of arenes (limited to
nitration, alkylation, acylation and chlorination);
describe nucleophilic
addition to carbonyl group
typified by the reactions with aldehydes and ketones of
HCN, hydroxylamine and the hydrazines;
describe electrophilic addition to carbon-carbon double
bond as typified by reaction with alkenes of halogens and
hydrogen halides;
Markovnikov’s Rule
explain the basis of Markovnikov’s Rule in terms of the
stability of carbocations;
Homolytic mechanisms
describe the homolytic reactions involved in the
halogenation of alkanes and methylbenzene and in the
polymerisation of alkenes.
14.11
Organic laboratory preparations
Students are expected to be familiar with detailed
laboratory preparation and isolation of pure samples of
the following organic substances:
(a) a bromoalkane from the corresponding alcohol;
(b) nitrobenzene from benzene;
(c) ethanal from ethanol;
(d) an ester from the reaction of a carboxylic acid and an
alcohol;
(e) phenylamine from nitrobenzene.
36
AM Syllabus (2015): Chemistry
15
Practical Examination
15.1
The examination will primarily attempt to test
practical skills.
The practical examination will seek to test the
ability of candidates to:
(a)
(b)
(c)
15.2
manipulate chemicals and simple
apparatus in quantitative and qualitative
exercises;
observe
and
record
results
of
experimental work;
interpret these observations and deduce
correct inferences and conclusions based
both on qualitative and quantitative data.
Quantitative
exercises
including
the
measurement of mass, volume, temperature and
time may be set.
The following may be set:
volumetric analysis involving acids and alkalis,
Indicator colour changes will be specified except
for colour changes involving phenolphthalein
and methyl orange.
redox titrations,
Redox titrations may involve iodimetry
(experiments involving the formation of iodine
from the reaction of iodide with an oxidizing
agent followed by titration with thiosulfate) or
titration involving permanganate and a reducing
agent.
complexometric titration involving EDTA.
In such titrations, the colour change of the
indicator will be specified.
Titrations with silver nitrate will not be set.
15.3
Qualitative exercises involving observations of
reactions and requiring deductions on the
chemical nature of the substances will also be
set. These will involve both organic and
inorganic materials.
A knowledge of classical systematic analysis
will not be required and will not be tested.
It will be assumed that candidates are familiar
with the following species from which unknown
compounds will be set for qualitative analysis:
Binary mixtures involving salts, oxides and
metals may be set for qualitative analysis.
(i) ammonium, potassium, sodium, silver,
calcium, calcium, magnesium, strontium,
barium,
aluminium,
lead(II),
zinc,
chromium(III), manganese(II), manganese(IV),
iron(II),
iron(III),
cobalt(II),
nickel(II),
copper(I), copper(II);
Students should be able to test for and recognize
by physical or chemical means the following
gases: ammonia, carbon dioxide, carbon
monoxide, hydrogen, oxygen, sulfur dioxide,
hydrogen halides (HX, X = Cl, Br, I), the
halogens (Cl, Br, I), nitrogen dioxide, sulfur
dioxide, hydrogen sulfide and nitrogen (by
elimination).
(ii)
oxide,
hydroxide,
carbonate,
hydrogencarbonate,
nitrate(V),
nitrate(III),
carboxylate, sulfate(IV), sulfate(VI), thiosulfate,
chloride, bromide, iodide, chromate(VI),
dichromate(VI), phosphate, manganate(VII);
Students should be aware of the toxicity of a
number of these species.
(iii) the more common metals (e.g. Zn, Al, Cu,
Fe)
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AM Syllabus (2015): Chemistry
Organic substances containing the following
functional groups may be set as unknown
compounds in qualitative analysis:
Double and triple bonds (alkenes, alkynes),
hydroxyl (alcohols and phenols), carbonyl
(aldehydes and ketones) carboxyl, carboxyl
derivatives (carboxylate salts, esters, amides,
nitriles) and amino.
15.4
Candidates are instructed to use labcoats and to
wear eye protection (safety goggles) during the
practical examination and to avoid wearing
clothes or articles that increase the risk of
accident in the laboratory.
Both aliphatic and aromatic compounds may be
set.
A combination of functional groups may be set,
e.g. amino acids, hydroxyacids etc.
Candidates will not be allowed into the
examination room unless equipped with proper
safety gear.
Recommended Texts
Any of the following are suitable textbooks for the syllabus; the list is not comprehensive.
1.
Lister, T and Renshaw, J, Understanding chemistry for Advanced Level, 4th Edition, Nelson Thornes
Ltd., 1999.
2.
Ramsden, EN, A-Level Chemistry, 4th Edition, Nelson Thornes Ltd., 2000.
3.
Jones, L and Atkins, P, Chemistry: molecules, matter and change, 4th Edition, W H Freeman, 1999.
4.
Hill, G and Holman, J, Chemistry in context, 6th Edition, Nelson Thornes Ltd., 2011.
38